CALCIUM BINDING BY CITRATE

MeThe citrate molecule in urine is thought to protect against formation of calcium stones. This thought began as reasoning from chemistry, and culminated in clinical trials which substantiate the idea. As a result manufacturers produce citrate products for medicinal use, and doctors prescribe the medicine.

All this is a wonderful success story, a kind of perfection of the paradigm of translational science: From science to a treatment for patients that reduces illness from kidney stone disease.

But what, exactly, is the science? Can scientists not enjoy the story of such a success, physicians derive from it a deeper understanding of the drug they so regularly dispense and patients the comfort that a perfected knowledge support the rightness of their prescribed treatment?

Citrate

The Molecule

As usual in such diagrams the carbon atoms are simply angles or kinks. Reading from left to right, there is a carbon atom bound (the solid lines) to two oxygen atoms (‘O’), one with a single and the other with a double bond. These bonds represent sharing of electrons by the atoms.
The single bonded oxygen has an extra electron in its outer shell, so it carries a negative charge (-). Calcium is an atom with two positive charges, so the idea of calcium and citrate binding to each other comes naturally as one thinks about opposites attracting one another.

Next in line is another carbon atom; the kink means the carbon is linked to the carbons on its left and right and to two hydrogen atoms. The hydrogen atoms are necessary because every carbon atom makes four bonds.

After that, is the third carbon which is very occupied. It has an oxygen which is itself bound to a hydrogen – a hydroxyl molecule, really 2/3 of a water molecule – and another carbon bound to two oxygens, one of which has a negative charge. To the right of this busy carbon the molecule repeats itself as in a mirror.

How beautiful nature is, how powerful its symmetries and suggestive its forms!

It is as though some great sculptress were taken with an image of perfection that a string of carbons might take the perfect form to mate with calcium, tiny in comparison, and doubly charged positive.

But how? How would the mating occur? If you do not look down, could you have imagined it?

Binding Constants

Calculation and experimental determination of calcium binding by citrate is complex. Partly, all 3 oxygens can accept a proton, so CALCIUM ION VS CITRATE MOLARITYthe acidity of the solution – urine in our case – matters. Partly, binding is complex. As shown in the section below, it involves forming a ring structure and a bi-molecular structure. In general calculations are performed using computer programs.

But simple experiments give a reasonable gauge of the power of citrate to bind calcium. In the figure adapted from Table 2 of the reference at the top of this section, when the molarity of calcium and citrate are equal, at both 2.5 and 5 mmol/l of total calcium (2.5 – grey circles and 5 – black circles in the legend), a common range in urine, only about 1 mmol of calcium is unbound and therefore a free ion that can combine with oxalate or phosphate to make a stone. From the shapes of the graphs, citrate is a powerful binder as the calcium ion falls almost linearly with citrate molarity.

This graph is very approximate. Actual calculations of citrate binding effects have to consider pH, ionic strength, and many varieties of citrate calcium salts. These are part of how supersaturations are calculatedYet for all its simplifications, this graph of ancient data suffices to show what citrate can do as a protection against urine crystallizations of calcium salts like calcium oxalate and calcium phosphate. 

Calcium Citrate Crystal

But this is not a complete story. What if calcium and citrate combine to make a crystal which becomes yet another kind of stone? They can indeed form a crystal, but one which is so soluble it is never a stone risk. Even so, how the crystal forms is a way to show how the molecule binds calcium, which is in itself simply very interesting.

Creation of the Di-Citrate

I have made this structure the featured illustration for the article, but put a copy here for visual convenience. Would you have imagined the two ends of the molecule would bend around to hold the calcium, which makes what was linear into what is, now, a ring?

Just below is another ring, identical in character. Neither is a crystal, merely they are a pair of rings.

But, that odd outcropping of a carbon atom on the original molecule has its charged oxygen, and through another calcium, caught there, the two are linked.

Citrate has two ways to bind calcium. A single molecule can bind one calcium atom. Two citrate molecules together can bind one calcium atom. So the ratio is 1.5 calcium atoms per molecule of citrate (3 atoms/2 citrate molecules).

The Making of a Crystal

How these paired rings make a crystal is not an easy story to tell. Here is a good reference which I will explicate not as a crystallographer, which I am certainly not, but as a narrator telling a good story.

calcium citrate crystal jpeg versioinThe calcium citrate crystal is built up out of repeating units, each of which is a pair of citrate molecules linked by a calcium atom. This ‘di-citrate’ unit has three calcium atoms in it, one in the center, which is unique, and one at each end which mirror each other. These are in the diagram just above.

All three coordinate with 8 oxygen atoms. By this I mean that in addition to the 2 oxygens shown in the simple figure of the di-citrate molecule, oxygens are shared that belong to di-citrate molecules ‘above; and ‘below’  and to the sides so as to make a set of plates like the floors of a parking building.

This complex macrame shows all three calcium atoms. The middle one – calcium 1, at the center of this drawing – relates to 6 oxygens bound to carbon atoms, and to two belonging to water molecules that are pulled into the final structure. The second – at the left lower corner – is coordinated by 8 oxygens bound to carbons, and the third – at the right lower corner, is coordinated with 7 oxygens bound to carbon and one hydroxyl (OH) bound to a carbon (see the simple di-citrate drawing).

This crystal, calcium citrate, is actually a medicinal product (‘Citracal’™) which when swallowed in a pill form dissolves in the gastrointestinal tract to donate calcium and citrate that can be absorbed into the blood. The crystal itself is not found in urine. The medication is of no immediate interest concerning stone disease because one does not use it as a treatment for stones but rather as a supplement for bones. Whether or not it might be helpful in preventing stones would require a trial.

The citrate used for stones is potassium citrate, which is simply the single citrate molecule with its 3 negative charges satisfied by potassium ions or protons.

Why, then, have I troubled you with the elaborate business of citrate calcium binding and crystal formation?

Because it is one way that citrate protects against stones. The molecule binds calcium which is therefore no longer free to combine with oxalate or phosphate to form kidney stones. The crystal calcium and citrate forms does not make stones because it is very soluble. If it were not, citrate would not be a protection against stones but merely the substrate for yet another calcium type stone.

Solubility of Calcium Citrate

What, then, is the evidence for this statement – that the calcium citrate crystal is so soluble that it does not make stones? I have said it several times but have provided you with no proof.

The three oxygens of citrate are partially occupied by protons, and when calcium citrate solubulutythey are they are not available to coordinate with calcium to make the di-citrate and its crystal. Therefore the solubility of citrate will be influenced by the concentration of protons, the acidity of the urine, represented here by pH.

The experiment is done by adding crystals to a simple salt solution, letting them equilibrate with the solution at a constant temperature, and measuring, in this case, the concentration of calcium that is in the solution, having left the crystals as they dissolve.

Two of the three oxygens are ‘weak’ acids which are half saturated with protons at pH values of about 3 and 4, meaning that throughout the range of urine acidity – pH 4.5 to 8 – both are free to bind with calcium. The third is at pH 6.4. One might expect an increase in calcium binding as pH rises above this point, but there is no obvious change in the solubility of the crystal between 5 and above 7.

At pH of 6, the mean for normal urine, the concentration of calcium in solution is about 0.2 mg/ml or 200 mg/liter. Given the atomic weight of calcium – 40- this is 5 mmol/liter. Calcium oxalate crystals dissolved in the same way yield a calcium concentration below 0.005 mg/ml or 5 mg/liter which is about 0.05 mmol/liter. Calcium phosphate crystals give a value of 0.08 mg calcium/ml. 80 mg/liter or 2 mmol/l, less than half of the citrate.

Since calcium and citrate are released from the crystal in proportions of 1.5 calcium per di-citrate, one presumes the equilibrium citrate molarity will be 66% of calcium or 3.3 mmol. Given the molecular weight of citrate is 192 mg/mmol, this amounts to 633 mg/liter of citrate. That is a high concentration of citrate, given that common excretion rates are rarely above 750 mg/day and urine volumes about 1.5 liters a day. Even so, some urine samples almost certainly achieve these concentrations of calcium and citrate on occasion. But equilibrium is not enough to create new crystals; one needs to achieve a higher value so that new crystal nuclei will form. That will be very unlikely for calcium citrate.

So citrate can combine with calcium to remove it from binding with oxalate and phosphate, and form a crystal of considerable solubility. Being very soluble, calcium citrate is rarely if ever found as a kidney stone.

Calcium and Citrate in Urine

calcium - citrate picture

One of two crucial issues about citrate in stone prevention is the relationship between the concentration of calcium and that of citrate. The higher the concentration of citrate compared to calcium, the lower the concentration of unbound calcium, which is free to combine with oxalate or phosphate to make kidney stones.

This graph from our research work shows the difference between urine calcium and urine citrate concentrations for normal people (red), and calcium phosphate (blue) and calcium oxalate (yellow) stone formers over the full day and overnight periods. These two kinds of stones have already been reviewed on this site. 

Normal people have lower urine calcium excretion rates than stone forming patients, but about the same excretion rates for citrate, so the calcium – citrate difference is below 1 mmol/liter. In the graph above which shows free calcium ion at two concentrations of total calcium – 2.5 and 5 mmol/l vs. citrate concentration, when the calcium and citrate concentrations are equal the free calcium is below 2 mmol/liter. When citrate concentration exceeds that of calcium, the free calcium will be lower.

Both of the patient groups have much higher calcium excretions than normal people and because their urine citrate excretions are no higher, and perhaps even lower than among normals, the concentrations of free calcium are much higher, in the range of 2 – 3 mmol/liter.

It is this kind of information which has long made scientists believe that citrate is an important factor in the normal defense against calcium stone formation, and which led to the successful trials which proved that this believe is not unfounded.

What is the Real Science?

The citrate story illustrates all the three forms of scientific research.

Empirical science is what we would call the meticulous measurement of the binding constants between citrate and calcium, and the specific structure of the calcium citrate crystal. It is also what we would call the pretty graph of urine calcium – citrate differences in normal people and stone forming patients.

Applied science is the trials which showed that the intuition of citrate as a treatment was a true intuition. It is indeed a treatment, and that is a fact which time will not alter.

Basic science, however, is not so obvious here. Where in the story do we encounter the passion or curiosity to ask how citrate has come to be in urine.

Certainly citric acid plays a role in biology vastly  – one might say infinitely – greater than that of stone prevention. It is the key molecule in the citric acid cycle which is so well known that I have only to reference it from Wikipedia. Known to schoolboys and schoolgirls everywhere, this cycle is used by all aerobic organisms to generate energy, and is of an extreme ancient origin.

Surely a molecule of such lineage and power is ruled little if at all by the problems of renal crystals. Yet it is handled by the kidneys with considerable finesse as if somehow important to the renal system or – perhaps – as if the renal systems were somehow important in the larger matters of maintaining serum levels of the molecule.

Here is imaginative science. Here is the place where a question of underlying cause comes into focus. Here is where nature presents issues of monumental consequence.

 

 

 

 

26 Responses to “CALCIUM BINDING BY CITRATE”

  1. Anthony J Perrotta

    Dear Dr. Coe,

    The chelating of the divalent calcium ion by citrate anion to keep the calcium ion in solution rather than combining with oxalate or phosphate anions to form stones. Basically, as I understand it, this chelating effect prevents supersaturation of the combining of oxalate or phosphate with calcium to form stones.

    What fruit juice would you recommend to supply the citrate anion to give the required chelating effect?

    Anthony Perrotta (PhD Chicago, 1965)

    Reply
    • Fredric Coe, MD

      Hi Anthony, indeed citrate complexes with urine calcium to form a very soluble salt. That reduces free calcium molarity. Oxygens on citrate also bind to surface calcium ions in calcium oxalate and calcium phosphate crystal nuclei competing against oxygens from phosphate and oxalate. This tends to disrupt orderly crystal growth and can provoke dissolution by disordering the surface charge array. Urine citrate is determined by the NaDC1 transporter that responds to pH of the proximal tubule fluid, so any alkali will raise urine citrate. Fruit juices are a miserable way to provide metabolizable anions because they collect the fructose from many fruits into a concentrate. Fructose – all sugars – raise urine calcium and fructose is a genuine health hazard. The best is some form of potassium citrate or else a diverse array of fruits – not the juices – and veggies that all contain abundant Krebs cycle precursors. Best, Fred (MD Chicago, 1961)

      Reply
  2. Susan

    Hi Fred, I have calcific rotator cuff tendonitis. Does taking Potassium Citrate help break down calcifications in the shoulder at all? Apparently, I’m excreting electrolytes in my urine. My doctor first said it was because I was taking D-mannose, but I haven’t been taking D-mannose for months and my urine electrolytes are still high. My doctor doesn’t have an answer as to why. Do some people excrete more potassium in their urine than others, and if so, why is that? Could excreting more potassium causes the body to be more prone to forming calcifications? Thank you, Susan

    Reply
    • Fredric Coe, MD

      Hi Susan, No; it will not help calcified rotator cuff disease. Taking citrate will not influence blood citrate. Anything that lowered serum calcium ion so joint crystals dissolved would endanger life – the brain, heart, etc all require serum calcium levels be very constant. As to urine potassium it faithfully follows diet and drug potassium intake – it is an atom. Likewise for sodium. So, you must eat more potassium than is usual – fruits and veggies are the usual source. But be happy; on average our urine potassium runs 40 – 60 mEq/day. The US ideal is over 100 mEq/day. Potassium intake has not effect on joint calcium content. Regards, Fred Coe

      Reply

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